Environmental Geochemistry

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Environmental Geochemistry. January 26, 2007. What is geochemistry?. The study of -chemical composition of the Earth and other planets -chemical processes and reactions that govern the composition of rocks and soils
Environmental Geochemistry

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Environmental geochemistry l.jpgSlide 1

Environmental Geochemistry

January 26, 2007

What is geochemistry l.jpgSlide 2

What is geochemistry?

The study of

-chemical composition of the Earth and other planets

-chemical processes and reactions that govern the composition of rocks and soils

-the cycles of matter and energy that transport the Earth's chemical components in time and space

-and their interaction with the hydrosphere and the atmosphere.

Outline of topics l.jpgSlide 3

Outline of Topics

  • Formation of the elements

  • Composition of Earth

  • Aqueous Solutions

  • Chemical Equilibrium

  • Acid-Base Equilibria

  • Redox

  • Biogeochemistry

  • Stable Isotopes

    (with comments on weathering, sorption, pollution…)

Formation of the elements l.jpgSlide 4

Formation of the Elements

Composition of earth l.jpgSlide 5

Composition of Earth

Aqueous solutions l.jpgSlide 7

Aqueous Solutions

Water is special

Slide9 l.jpgSlide 9

Ionic Strength

I = 1/2 ∑mz2

Chemical equilibrium l.jpgSlide 10

Chemical Equilibrium

Exists when a system is in a state of minimum energy (G)

  • - Often not completely attained in nature (e.g., photosynthesis leaves products out of chemical equilibrium)

  • - A good approximation of real world

  • Gives direction in which changes can take place (in the absence of energy input.)

  • Systems, including biological systems, can only move toward equilibrium.

  • -Gives a rough approximation for calculating rates of processes because, in

  • general, the farther a system is from equilibrium, the more rapidly it will move

  • toward equilibrium; however, it is generally not possible to calculate reaction rates from thermodynamic data.

Reaction rates equilibrium l.jpgSlide 11

Reaction Rates/Equilibrium

Acid base equilibria l.jpgSlide 12





Acid-Base Equilibria

Bronsted-Lowry definition: acid donates H+; base accepts H+

In aqueous systems, all acids stronger that H2O

generate excess H+ ions (or H3O+); all bases stronger

than H2O generate excess OH-

Acid base l.jpgSlide 13


Many reactions influence pH

Photosynthesis and respiration are acid base reactions.

aCO2(g) + bNO3- + cHPO42- + dSO42- + f Na+ + gCa2+ + hMg2+ iK+ + mH2O + (b + 2c + 2d -f -2g - 2h - i)H+<-----> {CaNbPcSdNafCagMghKiH2Om}biomass + (a + 2b)O2

Oxidation reactions often produce acidity.

Reduction reactions consume acidity

pH influences many processes

-weathering (Fe and Al more soluble at lower pH)

-cation exchange (leaching of base cations from soil due to acid rain)

-sorption(influences surface charge on minerals and therefore what sticks to them)

Acid base14 l.jpgSlide 14


Alkalinity ≈ ANC

Alkalinity = ∑(base cations) - ∑(strong acid anions)

Any process that affects the balance between base cations and acid anions must affect alkalinity.

Redox l.jpgSlide 15


  • The oxidation state of an atom is defined with the following

  • convention:

  • The oxidation state of an atom in an elemental form is 0.

  • In O2, O is in the 0 oxidation state.

  • When bonded to something else, oxygen is in oxidation

  • state -2 and hydrogen is in oxidation state of +1 (except for

  • peroxide and superoxide).

  • In CO32-, O is in -2 state, C is in +4 state.

  • The oxidation state of a single-atom ion is the charge on

  • the ion.

  • For Fe2+, Fe is in +2 oxidation state.

Redox16 l.jpgSlide 16


Redox reactions tend to be slow and are often out of thermodynamic equilibrium - but life exploits redox disequilibrium.

Oxidation - lose electrons

Reduction - gain electrons

Fe was oxidized, Mn was reduced

Why do we care about redox rxns l.jpgSlide 17

Why do we care about redox rxns?

Oxidation state can impact

  • Sorption/desorption

  • Solubility

  • Toxicity

  • Biological uptake


Measure of oxidation-reduction potential gives us info about chemical species present and microbes we may find.

Accumulation of o 2 in the atmosphere l.jpgSlide 18

Accumulation of O2 in the Atmosphere

Fe2+ = Fe(II) = slightly soluble in sea water with no O2 present

Add O2 - oxidizes Fe(II)-->Fe(III)

Very small [O2] required

Fe3+ = Fe(III) = extremely insoluble in water

Essentially all of the oxygen in the atmosphere came from photosynthesis

Biogeochemistry l.jpgSlide 19


Slide21 l.jpgSlide 21


ammonia→ nitrite → nitrate


nitrate → nitrite → nitric oxide → nitrous oxide → N2

N Fixation

N2 →ammonia

What is an isotope l.jpgSlide 23

What is an isotope?

  • Isotope- line of equal Z. It has the same # protons (ie. they are the same element) but a diff. # of neutrons.








How did all this stuff get here l.jpgSlide 24

How did all this stuff get here?

  • 4 types of isotopes, based on how they formed:

    • Primordial (formed w/ the universe)

    • Cosmogenic (made in the atmosphere)

    • Anthropogenic (made in bombs, etc)

    • Radiogenic (formed as a decay product)

Stable isotopes l.jpgSlide 25

Stable Isotopes

Light isotopes are fractionated during chemical reactions, phase changes, and biological reactions, leading to geographical variations in their isotopic compositions

FRACTIONATION: separation between isotopes on the basis of mass (usually), fractionation factor depends on temperature

Bonds between heavier isotopes are harder to break

Stable isotope examples l.jpgSlide 26

Stable Isotope Examples

  • Rayleigh fractionation: light isotopes evaporate more easily, and heavy isotopes rain out more quickly

d= {(Rsample – Rstandard) / Rstandard} x 103

Stable isotope examples27 l.jpgSlide 27

Stable Isotope Examples

  • d18Ocarbonate in forams depends on d18Oseawater as well as T, S

  • d18Oseawater depends on how much glacial ice there is

    • Glacial ice is isotopically light b/c of Rayleigh fract.

    • More ice means higher


Stable isotopes28 l.jpgSlide 28

Stable Isotopes

  • C inorganic matter, fossil fuels, and hydrocarbon gases is depleted in 13C ==> photosynthesis

    • used as an indicator of their biogenic origin and as a sign for the existence of life in Early Archean time (~ 3.8 billion years ago)

  • N isotopic composition of groundwater strongly affected by isotope fractionation in soils plus agricultural activities (use of N-fertilizer and discharge of animal waste)

  • Particulate matter in ocean enriched in 15N by oxidative degradation as particles sink through water column

    • Used for mixing and sedimentation studies

  • S isotopes fractionated during reduction of SO42- to S2- by bacteria

    • didn’t become important until after ~2.35 Ga when photosynthetic S-oxidizing bacteria had increased sulfate concentration in the oceans sufficiently for anaerobic S-reducing bacteria to evolve (photosynthesis preceded S-reduction which was followed by O respiration)

Stable isotope examples29 l.jpgSlide 29

Stable Isotope Examples

  • Stable isotopes can also tell you about biology

  • Organisms take up light isotopes preferentially

  • So, when an organism has higher

    30Si, it means that it was feeding from a depleted nutrient pool

Stable isotopes30 l.jpgSlide 30

Stable Isotopes

  • Boron isotopes measured in forams used for paleo-pH

    • d11B depends on pH

    • (Gary Hemming)

  • Nitrogen isotopes used for rapid temp. changes in ice cores

    • d15N depends on temp. gradient in firn

    • (Jeff Severinghaus)

  • Stable isotopes are also used to study magmatic processes, water-rock interactions, biological processes and anthropology and various aspects of paleoclimate

References l.jpgSlide 31


(WM White Geochemistry Ch9 - Stable Isotopes)

Isotopes: Principles and Applications - Faure & Mensing

How to Build a Habitable Planet - Wally Broecker

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